One of the key factors in any corrosion situation is the

1
Corrosion Fundamentals

• One of the key factors in any corrosion situation
  is the
• environment.
• The definition and characteristics of this
  variable can be quite
• complex.
• For practical, it is important to realize that
  the environment is
• a variable that can change with time and
  conditions.
• Environment that actually affects a metal
  corresponds to the
• microenvironmental conditions that this metal
  really sees, i.e.,
• the local environment at the surface of the
  metal.
• It is indeed the reactivity of this local
  environment that will
• determine the real corrosion damage.
• Thus, an experiment that investigates only the
  nominal
• environmental condition without consideration
  of local effects
• is useless for lifetime prediction.

2
water

• In our societies, water is used for a wide
  variety of purposes, from supporting life as
  potable water to performing a multitude of
  industrial tasks such as heat exchange and waste
  transport.
• The impact of water on the integrity of materials
  is thus an important aspect of system management.
  Since steels and other iron-based alloys are the
  metallic materials most commonly exposed to
  water, aqueous corrosion will be discussed with a
  special focus on the reactions of iron (Fe) with
  water (H2O).
• Metal ions go into solution at anodic areas in an
  amount chemically equivalent to the reaction at
  cathodic areas (Fig. 1.1).

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Reaction
4
Anodic Reaction

• In the cases of iron-based alloys, the following
  reaction usually takes place at anodic areas

5
Cathodic Reaction

• When iron corrodes, the rate is usually
  controlled by the cathodic reaction, which in
  general is much slower (cathodic control).
• In de-aerated solutions, the cathodic reaction is

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• The cathodic reaction proceeds rapidly in acids,
  but only slowly in alkaline or neutral aqueous
  media. The corrosion rate of iron in deaerated
  neutral water at room temperature, for example,
  is less than 5 µm/year. The rate of hydrogen
  evolution at a specific pH depends on the
  presence or absence of low-hydrogen overvoltage
  impurities in the metal.

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• The cathodic reaction can be accelerated by the
  reduction of dissolved oxygen in accordance with
  the following reaction, a process called
  depolarization

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• Dissolved oxygen reacts with hydrogen atoms
  adsorbed at random on the iron surface,
  independent of the presence or absence of
  impurities in the metal. The oxidation reaction
  proceeds as rapidly as oxygen reaches the metal
  surface. Adding (1.1) and (1.3), making use of
  the reaction H2O ??H OH-, leads to reaction
  (1.4),

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Diffusion-barrier Layer

• Hydrous ferrous oxide (FeO . nH2O) or ferrous
  hydroxide Fe(OH)2 composes the
  diffusion-barrier layer next to the iron surface
  through which O2 must diffuse. The pH of a
  saturated Fe(OH)2 solution is about 9.5, so that
  the surface of iron corroding in aerated pure
  water is always alkaline. The color of Fe(OH)2,
  although white when the substance is pure, is
  normally green to greenish black because of
  incipient oxidation by air.

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• At the outer surface of the oxide film, access to
  dissolved oxygen converts ferrous oxide to
  hydrous ferric oxide or ferric hydroxide, in
  accordance with

Hydrous ferric oxide is orange to red-brown in
color and makes up most of ordinary rust. It
exists as nonmagnetic ?Fe2O3 (hematite) or as
magnetic ?Fe2O3, the ? form having the greater
negative free energy of formation (greater
thermodynamic stability). Saturated Fe(OH)3 is
nearly neutral in pH. A magnetic hydrous ferrous
ferrite, Fe3O4 . nH2O, often forms a black
intermediate layer between hydrous Fe2O3 and FeO.
Hence rust films normally consist of three
layers of iron oxides in different states of
oxidation.
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Applications of Potential-pH Diagrams

• E-pH or Pourbaix diagrams are a convenient way of
  summarizing much thermodynamic data and provide a
  useful means of summarizing the thermodynamic
  behavior of a metal and associated species in
  given environmental conditions. E-pH diagrams are
  typically plotted for various equilibria on
  normal cartesian coordinates with potential (E)
  as the ordinate (y axis) and pH as the abscissa
  (x axis).

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Applications of Potential-pH Diagrams

• For corrosion in aqueous media, two fundamental
  variables, namely corrosion potential and pH, are
  deemed to be particularly important.
• Changes in other variables, such as the oxygen
  concentration, tend to be reflected by changes in
  the corrosion potential.
• Considering these two fundamental parameters,
  Staehle introduced the concept of overlapping
  mode definition and environmental definition
  diagrams, to determine under what environmental
  circumstances a given mode/submode of corrosion
  damage could occur (Fig. 1.2).

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Overlapping Modes
14
Applications of Potential-pH Diagrams

• In the application of E-pH diagrams to corrosion,
  thermodynamic data can be used to map out the
  occurrence of corrosion, passivity, and nobility
  of a metal as a function of pH and potential. The
  operating environment can also be specified with
  the same coordinates, facilitating a
  thermodynamic prediction of the nature of
  corrosion damage.
• A particular environmental diagram showing the
  thermodynamic stability of different chemical
  species associated with water can also be derived
  thermodynamically. This diagram, which can be
  conveniently superimposed on E-pH diagrams, is
  shown in Fig. 1.3. While the E-pH diagram
  provides no kinetic information whatsoever, it
  defines the thermodynamic boundaries for
  important corrosion species and reaction

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Figure 1.3 Thermodynamic stability of water,
oxygen, and hydrogen. (A is the Equilibrium line
for the reaction H2 2H 2e-. B is the
equilibrium line for the reaction 2H2O O2
4H 4e-. indicates increasing thermodynamic
driving force for cathodic oxygen reduction,
as the potential falls below line B.
indicates increasing thermodynamic driving force
for cathodic hydrogen evolution, as the
potential falls below line A.)
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Observed Corrosion

• The observed corrosion behavior of a particular
  metal or alloy can also be superimposed on E-pH
  diagrams. Such a superposition is presented in
  Fig. 1.4. The corrosion behavior of steel
  presented in this figure was characterized at
  different potentials in solutions with varying pH
  levels.

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Thermodynamic boundaries of the types of
corrosion observed on steel
18
Corrosion of steel in water at elevated
temperatures

• Many phenomena associated with corrosion damage
  to iron-based alloys in water at elevated
  temperatures can be rationalized on the basis of
  iron-water E-pH diagrams. Marine boilers on ships
  and hot-water heating systems for buildings are
  relevant practical examples

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Marine boilers

• Two important variables affecting water-side
  corrosion of iron-based alloys in marine boilers
  are the pH and oxygen content of the water.
• As the oxygen level has a strong influence on the
  corrosion potential, these two variables exert a
  direct influence in defining the position on the
  E pH diagram. A higher degree of aeration raises
  the corrosion potential of iron in water, while a
  lower oxygen content reduces it.

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Elevated-temperature and Ambient-temperature

• When considering the water-side corrosion of
  steel in marine boilers, both the
  elevated-temperature and ambient-temperature
  cases should be considered, since the latter is
  important during shutdown periods. Boiler
  feedwater treatment is an important element of
  minimizing corrosion damage

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• A fundamental treatment requirement is
  maintaining an alkaline pH value, ideally in the
  range of 10.5 to 11 at room temperature. This
  precaution takes the active corrosion field on
  the left-hand side of the E-pH diagrams out of
  play, as shown in the E-pH diagrams drawn for
  steel at two temperatures, 25C (Fig. 1.5) and
  210C (Fig. 1.6).

22
E-pH diagram of iron in water at 25C and its
observed corrosion behavior
23
Corrosion forms of Steel in Water
24
E-pH diagram of iron in water at 210C
25
Observations of Figure 1.5

• At the recommended pH levels, around 11, the E-pH
  diagram in Fig. 1.5 indicates the presence of
  thermodynamically stable oxides above the zone of
  immunity.
• It is the presence of these oxides on the surface
  that protects steel from corrosion damage in
  boilers.

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Practical Observations

• Practical Practical experience related to boiler
  corrosion kinetics at different feed water pH
  levels is included in Fig. 1.5. The kinetic
  information in Fig. 1.5 indicates that high
  oxygen contents are generally undesirable. It
  should also be noted from Figs. 1.5 and 1.6 that
  active corrosion is possible in acidified
  untreated boiler water, even in the absence of
  oxygen.

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Localized Pitting

• Inspection of the kinetic data presented in Fig.
  1.5 reveals a tendency for localized pitting
  corrosion at feed water pH levels between 6 and
  10. This pH range represents a situation in
  between complete surface coverage by protective
  oxide films and the absence of protective films.

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• Localized anodic dissolution is to be expected on
  a steel surface covered by a discontinuous oxide
  film, with the oxide film acting as a cathode.
• Another type of localized corrosion, caustic
  corrosion, can occur when the pH is raised
  excessively on a localized scale. The E-pH
  diagrams in Figs. 1.5 and 1.6 indicate the
  possibility of corrosion damage at the high end
  of the pH axis, where the protective oxides are
  no longer stable.
• Such undesirable pH excursions tend to occur in
  high temperature zones, where boiling has led to
  a localized caustic concentration.
• A further corrosion problem, which can arise in
  highly alkaline environments, is caustic
  cracking, a form of stress corrosion cracking.
  Examples in which such microenvironments have
  been proven include seams, rivets, and boiler
  tube-to-tube plate joints.

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Another Example

• Hydronic Heating of Buildings

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Figure 1.7 E-pH diagram of iron in water at
25C, highlighting the corrosion processes in the
hydronic pH range
31
Figure 1.8 E-pH diagram of iron in water at
85C (hydronic system
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Hydronic Heating of Buildings

• Given a pH range for mains water of 6.5 to 8 and
  the E-pH diagrams in Figs. 1.7 (25C) and 1.8
  (85C), it is apparent that minimal corrosion
  damage is to be expected if the corrosion
  potential remains below _0.65 V (SHE).
• The position of the oxygen reduction line
  indicates that the cathodic oxygen reduction
  reaction is thermodynamically very favorable.

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Role of Oxygen Content

• From kinetic considerations, the oxygen content
  will be an important factor in determining
  corrosion rates.
• The oxygen content of the water is usually
  minimal, since the solubility of oxygen in water
  decreases with increasing temperature (Fig. 1.9),
  and any oxygen remaining in the hot water is
  consumed over time by the cathodic corrosion
  reaction.
• Typically, oxygen concentrations stabilize at
  very low levels (around 0.3 ppm), where the
  cathodic oxygen reduction reaction is stifled and
  further corrosion is negligible.

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Figure 1.9 Solubility of oxygen in water in
equilibrium with air at different temperatures.
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Role of Oxygen Content

• Higher oxygen levels in the system drastically
  change the situation, potentially reducing
  radiator lifetimes by a factor of 15.
• The undesirable oxygen pickup is possible during
  repairs, from additions of fresh water to
  compensate for evaporation, or, importantly,
  through design faults that lead to continual
  oxygen pickup from the expansion tank.
• The higher oxygen concentration shifts the
  corrosion potential to higher values, as shown in
  Fig. 1.7. Since the Fe(OH)3 field comes into play
  at these high potential values, the accumulation
  of a red-brown sludge in radiators is evidence of
  oxygen contamination.

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Hydrogen Production

• From the E-pH diagrams in Figs. 1.7 and 1.8, it
  is apparent that for a given corrosion potential,
  the hydrogen production is thermodynamically more
  favorable at low pH values.
• The production of hydrogen is, in fact, quite
  common in microenvironments where the pH can be
  lowered to very low values, leading to severe
  corrosion damage even at very low oxygen levels.
• The corrosive microenvironment prevailing under
  surface deposits is very different from the bulk
  solution. In particular, the pH of such
  microenvironments tends to be very acidic.

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